Ionic Theory and Electrolytes in Solution

Ionic Theory of Solutions

• Electrolytes dissolve in water and produce an electrically conducting solution (e.g., salt solution).
• Nonelectrolytes dissolve in water but yield a non-conducting or poorly conducting solution (e.g., sucrose solution).
• Strong electrolytes dissolve completely in water to produce ions, allowing electrical conduction, such as lighting up a light bulb.

Strong and Weak Electrolytes

• Strong electrolytes exist almost entirely as ions in solution.
• Weak electrolytes dissolve in water and result in a relatively small percentage of ions; as a result, they only weakly conduct electricity.

Solubility of Compounds

• Soluble compounds dissolve readily in a solvent.
• Insoluble compounds dissolve very little in a solvent.

Solubility Rules

• Group IA and ammonium compounds are generally soluble.
• Acetates and nitrates are generally soluble.
• Most chlorides, bromides, and iodides are soluble, with exceptions like AgCl, Hg₂Cl₂, PbCl₂, AgBr, Hg₂Br₂, HgBr₂, PbBr₂, AgI, Hg₂I₂, HgI₂, and PbI₂.
• Most sulfates are soluble, with exceptions like CaSO₄, SrSO₄, BaSO₄, Ag₂SO₄, Hg₂SO₄, and PbSO₄.
• Most carbonates are insoluble, except those of Group IA and (NH₄)₂CO₃.
• Most phosphates are insoluble, except those of Group IA and (NH₄)₃PO₄.
• Most sulfides are insoluble, except those of Group IA and (NH₄)₂S.
• Most hydroxides are insoluble, except those of Group IA, Ca(OH)₂, Sr(OH)₂, and Ba(OH)₂.

Molecular and Ionic Equations

• Molecular Equations represent reactants and products as molecular substances, even if they exist as ions in solution; state symbols included are (s), (l), (g), (aq)
• Aqueous solutions of AgNO₃, NaCl, and NaNO₃ are written as compounds in the initial molecular equation.
• Example: AgNO3(aq) + NaCl(aq) → AgCl(s) + NaNO3(aq)

Complete Ionic Equations

• Strong electrolytes are written as separate ions in the solution.
• Other reactants and products are written in molecular form.
• State symbols include (s), (l), (g), and (aq).
• AgNO3(aq) + NaCl(aq) → AgCl(s) + NaNO3(aq) becomes Ag⁺(aq) + NO₃⁻(aq) + Na⁺(aq)Cl⁻(aq) → AgCl(s) + Na⁺(aq) + NO₃⁻(aq).

Spectator Ions and Net Ionic Equations

• Spectator ions are ions that do not participate in a reaction and appear on both sides of the ionic equation.
• Net ionic equations show only the species that participate in the reaction, with spectator ions omitted.
• For Ag+(aq) + NO3-(aq) + Na+(aq) + Cl-(aq) → AgCl(s) + Na+(aq) + NO3-(aq), the net ionic form is Ag+(aq) + Cl-(aq) → AgCl(s).

Determining Reaction Occurrence and Equation Types

• To determine if a reaction occurs, write the molecular, ionic, and net ionic equations.
• Start with a balanced molecular equation, predict possible products, and determine their solubility.
• If all potential products are soluble, there is no net reaction.

Molecular, Ionic, and Net Ionic Equation Example

• Molecular Equation: 2NaOH(aq) + MgCl₂(aq) → 2NaCl(aq) + Mg(OH)₂(s)
• Ionic Equation: 2Na⁺(aq) + 2OH⁻(aq) + Mg²⁺(aq) + 2Cl⁻(aq) → 2Na⁺(aq) + 2Cl⁻(aq) + Mg(OH)₂(s)
• Net Ionic Equation: 2OH⁻(aq) + Mg²⁺(aq) → Mg(OH)₂(s)

Precipitation Reactions

• Definition: Reactions that produce a solid ionic substance from mixing two solutions.
• A precipitate refers to the insoluble solid formed during a chemical reaction within a solution.
• Steps to Predict Precipitation Reactions:
  • Predict products by exchanging the ions.
  • Determine the state of each product as (s), (l), (g), or (aq).
  • Note: If all products are aqueous, there is no net reaction.

Acid Base Reactions

• Involve the transfer of a proton (H⁺) between reactants.

Arrhenius Definitions

• Arrhenius Acid: A substance producing hydrogen ions (H⁺) when dissolved in water.
• Arrhenius Base: A substance producing hydroxide ions (OH⁻) when dissolved in water.

Bronsted-Lowry Definitions

• Brønsted-Lowry Acid: A species donating a proton in a proton-transfer reaction.
• Brønsted-Lowry Base: A species accepting a proton in a proton-transfer reaction.

Acid-Base Indicator

• A dye used to distinguish between acidic and basic solutions based on color changes.

Strong and Weak Acids

• Strong Acid: Ionizes completely in water, exists entirely as ions, and is a strong electrolyte.
  • Common examples: HNO₃, H₂SO₄, HClO₄, and HCl.
• Weak Acid: Only partially ionizes in water, exists primarily as molecules with some ions, and is a weak electrolyte.
  • Example: Acetic acid (CH₃COOH).
• The strength of an acid is independent of its concentration.

Strong and Weak Bases

• Strong Base: Ionizes completely in water, exists entirely as ions, and is a strong electrolyte.
  • Examples of strong bases: LiOH, NaOH, KOH, and Ca(OH)₂.
• Weak Base: Only partly ionized in water, exists primarily as molecules with some ions, and is a weak electrolyte.
  • Often nitrogen bases such as NH3: NH₃(aq) + H₂O(l) ⇌ NH₄⁺(aq) + OH⁻(aq).

Classifying Acids and Bases

• KOH: strong base
• H₂S: weak acid
• CH₃NH₂: weak base
• HClO₄: strong acid

Neutralization Reaction

• Involves the reaction of an acid and a base to form an ionic compound (salt) and possibly water.
• H₂SO₃(aq) + 2KOH(aq) → 2H₂O(l) + K₂SO₃(aq) represents the balanced chemical equation for the neutralization of sulfurous acid by potassium hydroxide.

Acid Base Reaction with Gas Formation

• Some salts, such as sulfides, sulfites, and carbonates, when treated with an acid produce specific gases, such as hydrogen sulfide, sulfur trioxide, and carbon dioxide.

Examples of Gas-Formation in Acid-Base Reactions

• Na₂S(aq) + 2HCl(aq) → 2NaCl(aq) + H₂S(g)
• Na₂CO₃(aq) + 2HCl(aq) → 2NaCl(aq) + H₂O(l) + CO₂(g)
• Na₂SO₃(aq) + 2HCl(aq) → 2NaCl(aq) + H₂O(l) + SO₂(g)

Oxidation-Reduction (Redox) Reactions

• Redox reactions lead to an increase and decrease of oxidation numbers of the respective atoms in the reaction.
• Oxidation number is the hypothetical charge assigned to an atom, representing the total number of electrons gained or lost to form a chemical bond.
• Example: Mg has an oxidation number of +2 and Cl is -1 in MgCl₂.

Rules for Assigning Oxidation Numbers

• Oxidation number of an atom in an element is zero.
• Oxidation number of a monoatomic ion equals the charge on the ion (Na⁺ is +1).
• Oxidation number of oxygen is -2 in most compounds, but -1 in peroxides like H₂O₂.
• Oxidation number of hydrogen is +1 in most compounds, but -1 in binary compounds with a metal like CaH₂.
• Oxidation number of fluorine is -1. Other halogens can be -1 in binary compounds except when bonded to a halogen above them or to oxygen.
• The sum of oxidation numbers in a compound is zero; in a polyatomic ion, it equals the charge on the ion.

Oxidation Numbers: Example Calculation

• Oxidation number of Mn in KMnO4 is +7 [1(+1) + 1(x) + 4(-2) = 0].
• Oxidation number of Mn in K₂MnO₄ is +6 [2(+1) + 1(x) + 4(-2) = 0].

Oxidation-Reduction Half Reactions

• Oxidation: Loss of electrons by a species, indicated by an increase in oxidation number.
• Reduction: Gain of electrons by a species, indicated by a decrease in oxidation number.

Common Redox Reactions

• Combination
• Decomposition
• Displacement
• Combustion

Combination Reaction

• Two substances combine into a third substance.
• Example: 2Na(s) + Cl₂(g) → 2NaCl(s).

Decomposition Reaction

• A single compound reacts to produce two or more substances.
• Example: 2HgO(s) → 2Hg(l) + O₂(g)

Displacement Reaction

• An element reacts with a compound and displaces another element from it.
• Example: Zn(s) + 2HCl(aq) → H₂(g) + ZnCl₂(aq)

Combustion Reaction

• A substance reacts with oxygen, quickly releasing heat to produce a flame.
• Example: 4Fe(s) + 3O₂(g) → 2Fe₂O₃(s)

Balancing Redox Reactions (Half-Reaction Method)

• Identify oxidation numbers to determine what is oxidized and reduced.
• For the reaction Zn(s) + Ag⁺(aq) → Zn²⁺(aq) + Ag(s), Zn is oxidized (0 to +2) and Ag⁺ is reduced (+1 to 0).
• Write unbalanced half-reactions: Zn(s) → Zn²⁺(aq) (oxidation) and Ag⁺(aq) → Ag(s) (reduction).
• Balance charges by adding electrons: Zn(s) → Zn²⁺(aq) + 2e⁻ and e⁻ + Ag⁺(aq) → Ag(s).
• Multiply half-reactions to equalize electron count: 2e⁻ + 2Ag⁺(aq) → 2Ag(s).
• Add the half-reactions: Zn(s) + 2Ag⁺(aq) → Zn²⁺(aq) + 2Ag(s).

Learn about electrolytes, nonelectrolytes, and the differences between strong and weak electrolytes. Explore the concept of solubility and delve into solubility rules for various ionic compounds. Understand which compounds dissolve readily in a solvent and which remain insoluble.